An iron nail with a mass of 0.750 g was dissolved in dilute sulfuric acid and diluted to 100.0 mL volume. A 10.0 mL aliquot of this solution was then titrated with standard 0.200 M potassium permanganate solution, according to the following equation: MnO4- + H+ +Fe2+ --> Mn2+ + Fe3+ +H2O
It took 12.08 mL of the permanganate solution to oxidize all the iron.
a. Determine the mass of iron present in the original sample.
b. Given the result from (a), what was the purity of the iron nail?
c. If the nail contained a non-iron impurity that did not react with the sulfuric acid, but was oxidized by the permanganate, would the calculated purity of the nail be overestimated or underestimated?
1. First, balance the Fe/MnO4 equation but correct it first. The part that matters is 5Fe^2+ + MnO4^- ==> 5Fe^3+ + Mn^2+.
a. Millimols MnO4 used = M x mL = 12.08 x 0.200 = approx 2.4.
Using the coefficients in the balanced equation, convert mmols MnO4 to mmols Fe.
That's approx 2.4 x (5 Fe/1 MnO4) = approx 12 mmols Fe = 0.012 mols.
Convert to g Fe. Approx 0.012 x 55.85 = approx 0.7 g Fe in the 10.0 mL aliquot titrated or approx 7 g Fe in the origininal sample.
Note that this gives you a larger amount of Fe than the initial sample and that isn't possible. My best guess is that the MnO4 solution is 0.02 M and not 0.2 M. Look over the numbers and make sure they are right.
b. %purity = (mass Fe/mass sample)*100 ?
To solve this problem, we need to use the information provided and perform a series of calculations.
a. Determine the mass of iron present in the original sample:
We know that it took 12.08 mL of the 0.200 M potassium permanganate solution to react with the iron in the sample. From the balanced equation, we see that the ratio of the iron (Fe2+) to the potassium permanganate (MnO4-) is 1:1. This means that 1 mole of iron reacts with 1 mole of permanganate ion.
First, calculate the number of moles of potassium permanganate that reacted:
moles of KMnO4 = 0.200 moles/L * 0.01208 L = 0.002416 moles
Since the ratio is 1:1, we have the same number of moles of iron:
moles of iron = 0.002416 moles
Now, we can calculate the mass of iron using its molar mass:
mass of iron = moles of iron * molar mass of iron
The molar mass of iron (Fe) is 55.85 g/mol.
mass of iron = 0.002416 moles * 55.85 g/mol = 0.1349 g
Therefore, the mass of iron present in the original sample is approximately 0.1349 g.
b. Given the result from (a), what was the purity of the iron nail:
To determine the purity of the iron nail, we compare the mass of the iron in the sample (0.1349 g) to the mass of the entire sample (0.750 g).
purity of iron = (mass of iron / mass of sample) * 100%
purity of iron = (0.1349 g / 0.750 g) * 100% = 17.99%
Therefore, the purity of the iron nail is approximately 17.99%.
c. If the nail contained a non-iron impurity that did not react with the sulfuric acid, but was oxidized by the permanganate, would the calculated purity of the nail be overestimated or underestimated:
The calculated purity of the iron nail would be overestimated. This is because any non-iron impurity that reacts with the permanganate will also contribute to the total mass measured during the titration, leading to a higher calculated mass of iron. Therefore, the presence of a non-iron impurity that reacts with the permanganate would result in an overestimation of the purity of the iron nail.